Abundant supplies of fresh water are essential to the development of industry.
Enormous quantities are required for the cooling of products and equipment,
for process needs, for boiler feed, and for sanitary and potable water
supply.

THE PLANETARY WATER CYCLE

Industry is a small participant in the global water cycle .The finite amount of water on the planet participates in a very complicated recycling scheme that provides for its reuse. This recycling of water is termed the "Hydrologic Cycle" (see Figure 1-1).

Evaporation under the influence of sunlight takes water from a liquid to a
gaseous phase. The water may condense in clouds as the temperature drops
in the upper atmosphere. Wind transports the water over great distances
before releasing it in some form of precipitation. As the water condenses
and falls to the ground, it absorbs gases from the environment. This is
the principal cause of acid rain and acid snow.

WATER AS A SOLVENT

Pure water (H20) is colorless, tasteless,
and odorless. It is composed of hydrogen and oxygen. Because water becomes
contaminated by the substances with which it comes into contact, it is
not available for use in its pure state. To some degree, water can dissolve
every naturally occurring substance on the earth. Because of this property,
water has been termed a "universal solvent." Although beneficial to mankind,
the solvency power of water can pose a major threat to industrial equipment.
Corrosion reactions cause the slow dissolution of metals by water. Deposition
reactions, which produce scale on heat transfer surfaces, represent a
change in the solvency power of water as its temperature is varied. The
control of corrosion and scale is a major focus of water treatment technology.

WATER IMPURITIES

Water impurities include dissolved and suspended solids. Calcium bicarbonate
is a soluble salt. A solution of calcium bicarbonate is clear, because
the calcium and bicarbonate are present as atomic sized ions which are
not large enough to reflect light. Some soluble minerals impart a color
to the solution. Soluble iron salts produce pale yellow or green solutions;
some copper salts form intensely blue solutions. Although colored, these
solutions are clear. Suspended solids are substances that are not completely
soluble in water and are present as particles. These particles usually
impart a visible turbidity to the water. Dissolved and suspended solids
are present in most surface waters. Seawater is very high in soluble sodium
chloride; suspended sand and silt make it slightly cloudy. An extensive
list of soluble and suspended impurities found in water is given in Table
1-1.

refers to the sum of dissolved and suspended solids, determined gravimetrically

see "Dissolved Solids" and "Suspended Solids"

Surface Water

The ultimate course of rain or melting snow depends on the nature of
the terrain over which it flows. In areas consisting of hard packed clay,
very little water penetrates the ground. In these cases, the water generates
"runoff". The runoff collects in streams and rivers. The rivers empty
into bays and estuaries, and the water ultimately returns to the sea,
completing one major phase of the hydrologic cycle shown in Figure
1-1.

As water runs off along the surface, it stirs up and suspends particles of sand and soil, creating silt in the surface water. In addition, the streaming action erodes rocky surfaces, producing more sand. As the surface water cascades over rocks, it is aerated. The combination of oxygen, inorganic nutrients leached from the terrain, and sunlight supports a wide variety of life forms in the water, including algae, fungi, bacteria, small crustaceans, and fish.

Often, river beds are lined with trees, and drainage areas feeding the rivers are forested. Leaves and pine needles constitute a large percentage of the biological content of the water. After it dissolves in the water, this material becomes a major cause of fouling of ion exchange resin used in water treatment.

The physical and chemical characteristics of surface water contamination
vary considerably over time. A sudden storm can cause a dramatic short
term change in the composition of a water supply. Over a longer time period,
surface water chemistry varies with the seasons. During periods of high
rainfall, high runoff occurs. This can have a favorable or unfavorable
impact on the characteristics of the water, depending on the geochemistry
and biology of the terrain.

Surface water chemistry also varies over multi year or multidecade cycles
of drought and rainfall. Extended periods of drought severely affect the
availability of water for industrial use. Where rivers discharge into
the ocean, the incursion of salt water up the river during periods of
drought presents additional problems. Industrial users must take surface
water variability into account when designing water treatment plants and
programs.

Groundwater

Water that falls on porous terrains, such as sand or sandy loam, drains
or percolates into the ground. In these cases, the water encounters a
wide variety of mineral species arranged in complex layers, or strata.
The minerals may include granite, gneiss, basalt, and shale. In some cases,
there may be a layer of very permeable sand beneath impermeable clay.
Water often follows a complex three dimensional path in the ground. The
science of groundwater hydrology involves the tracking of these water
movements.

Table 1-2. A comparison of surface water and groundwater characteristics.

Characteristic

Surface Water

Ground Water

Turbidity

high

low

Dissolved minerals

low-moderate

high

Biological content

high

low

Temporal variability

very high

low

In contrast to surface supplies, groundwaters are relatively free from
suspended contaminants, because they are filtered as they move through
the strata. The filtration also removes most of the biological contamination.
Some groundwaters with a high iron content contain sulfate reducing bacteria.
These are a source of fouling and corrosion in industrial water systems.

Groundwater chemistry tends to be very stable over time. A groundwater
may contain an undesirable level of scale forming solids, but due to its
fairly consistent chemistry it may be treated effectively.

Mineral Reactions: As groundwater encounters different minerals, it dissolves
them according to their solubility characteristics. In some cases chemical
reactions occur, enhancing mineral solubility.

A good example is the reaction of groundwater with limestone. Water percolating
from the surface contains atmospheric gases. One of these gases is carbon
dioxide, which forms carbonic acid when dissolved in water. The decomposition
of organic matter beneath the surface is another source of carbon dioxide.
Limestone is a mixture of calcium and magnesium carbonate. The mineral,
which is basic, is only slightly soluble in neutral water. The slightly
acidic groundwater reacts with basic limestone in a neutralization reaction
that forms a salt and a water of neutralization. The salt formed by the
reaction is a mixture of calcium and magnesium bicarbonate. Both bicarbonates
are quite soluble. This reaction is the source of the most common deposition
and corrosion problems faced by industrial users. The calcium and magnesium
(hardness) form scale on heat transfer surfaces if the groundwater is
not treated before use in industrial cooling and boiler systems. In boiler
feedwater applications, the thermal breakdown of the bicarbonate in the
boiler leads to high levels of carbon dioxide in condensate return systems.
This can cause severe system corrosion.

Structurally, limestone is porous. That is, it contains small holes and
channels called "interstices". A large formation of limestone can hold
vast quantities of groundwater in its structure. Limestone formations
that contain these large quantities of water are called aquifers, a term
derived from Latin roots meaning water bearing.

If a well is drilled into a limestone aquifer, the water can he withdrawn
continuously for decades and used for domestic and industrial applications.
Unfortunately, the water is very hard, due to the neutralization/dissolution
reactions described above. This necessitates extensive water treatment
for most uses.

CHEMICAL REACTIONS

Numerous chemical tests must be conducted to ensure effective control of a water treatment program. Most of these tests are addressed in detail in Chapters 39-71. Because of their significance in many systems, three tests, pH, alkalinity, and silica, are discussed here as well.

pH Control

Good pH control is essential for effective control of deposition and corrosion
in many water systems. Therefore, it is important to have a good understanding
of the meaning of pH and the factors that affect it.

Pure H2O exists as an equilibrium between the acid species, H+ (more correctly expressed as a protonated water molecule, the hydronium ion, H30+) and the hydroxyl radical, OH -. In neutral water the acid concentration equals the hydroxyl concentration and at room temperature they both are present at 10-7 gram equivalents (or moles) per liter.

The "p" function is used in chemistry to handle very small numbers. It is the negative logarithm of the number being expressed. Water that has 10-7 gram equivalents per liter of hydrogen ions is said to have a pH of 7. Thus, a neutral solution exhibits a pH of 7. Table 1-3 lists the concentration of H+ over 14 orders of magnitude. As it varies, the concentration of OH - must also vary, but in the opposite direction, such that the product of the two remains constant.

Table 1-3. pH relationships.

pHa

H+ Concentration Exponential Notation, gram moles/L

H+ Concentration, Normality

OH - Concentration, Normality

OH - Concentration, Exponential Notation, gram moles/L

pOH -

0

100

1

0.00000000000001

10-14

14

1

10-1

0.1

0.0000000000001

10--13

13

2

10-2

0.01

0.000000000001

10--12

12

3

10-3

0.001

0.00000000001

10-11

11

4

10-4

0.0001

0.0000000001

10-10

10

5

10-5

0.00001

0.000000001

10-9

9

6

10-6

0.000001

0.00000001

10-8

8

7

10-7

0.0000001

0.0000001

10-7

7

8

10-8

0.00000001

0.000001

10-6

6

9

10-9

0.000000001

0.00001

10-5

5

10

10-10

0.0000000001

0.0001

10-4

4

11

10-11

0.00000000001

0.001

10-3

3

12

10-12

0.000000000001

0.01

10-2

2

13

10-13

0.0000000000001

0.1

10-1

1

14

10-14

0.00000000000001

1

100

0

apH+pOH=14.

Confusion regarding pH arises from two sources:

the inverse nature of the function

the pH meter scale

It is important to remember that as the acid concentration increases, the pH value decreases (see Table 1-4).

Table 1-4. Comparative pH levels of common solutions.

12

OH - alkalinity 500 ppm as CaCO3

11

OH - alkalinity 50 ppm as CaCO3
Columbus. OH, drinking water, a

10

OH - alkalinity 5 ppm as CaCO3

9

strong base anion exchanger effluents

8

phenolphthalein end point

7

neutral point at 25 °C

6

Weymouth, NIA, drinking water, a

5

methyl orange end point

4

FMA 4 ppm as CaCO3

3

FMA 40 ppm as CaCO3
strong acid cation exchanger effluent

2

FMA 400 ppm as CaCO3

a Extremes of drinking water pH

The pH meter can be a source of confusion, because the pH scale on the
meter is linear, extending from 0 to 14 in even increments. Because pH
is a logarithmic function, a change of I pH unit corresponds to a 10 fold
change in acid concentration. A decrease of 2 pH units represents a 100
fold change in acid concentration.

Alkalinity

Alkalinity tests are used to control lime-soda softening processes and boiler
blowdown and to predict the potential for calcium scaling in cooling water
systems. For most water systems, it is important to recognize the sources
of alkalinity and maintain proper alkalinity control.

Carbon dioxide dissolves in water as a gas. The dissolved carbon dioxide reacts with solvent
water molecules and forms carbonic acid according to the following reaction:

CO2 + H2O = H2CO3

Only a trace amount of carbonic acid is formed,
but it is acidic enough to lower pH from the
neutral point of 7. Carbonic acid is a weak acid,
so it does not lower pH below 4.3. However, this
level is low enough to cause significant corrosion
of system metals.

If the initial loading of CO2 is held constant
and the pH is raised, a gradual transformation into the bicarbonate ion HCO3-
occurs. This is shown in Figure
1-2.

The transformation is complete at pH 8.3. Further elevation of the pH
forces a second transformation into carbonate, CO32-.
The three species carbonic acid, bicarbonate, and carbonate can be converted
from one to another by means of changing the pH of the water.

Variations in pH can be reduced through "buffering" the addition of acid (or
caustic). When acid (or caustic) is added to a water containing carbonate/bicarbonate
species, the pH of the system does not change as quickly as it does in pure
water. Much of the added acid (or caustic) is consumed as the carbonate/bicarbonate
(or bicarbonate/carbonic acid) ratio is shifted.

Alkalinity is the ability of a natural water to neutralize acid (i.e., to reduce the pH depression expected from a strong acid by the buffering mechanism mentioned above). Confusion arises in that alkaline pH conditions exist at a pH above 7, whereas alkalinity in a natural water exists at a pH above 4.4.

Alkalinity is measured by a double titration; acid is added to a sample
to the Phenolphthalein end point (pH 8.3) and the Methyl Orange end point
(pH 4.4). Titration to the Phenolphthalein end point (the P-alkalinity)
measures OH - and 1/2 CO32-; titration
to the Methyl Orange end point (the M-alkalinity) measures OH -,
CO32-
and HCO3 .

Silica

When not properly controlled, silica forms highly insulating, difficult to remove deposits in cooling systems, boilers, and turbines. An understanding of some of the possible variations in silica testing is valuable.

Most salts, although present as complicated crystalline structures in the solid phase, assume fairly simple ionic forms in solution. Silica exhibits complicated structures even in solution.

Silica exists in a wide range of structures, from a simple silicate to a complicated polymeric material. The polymeric structure can persist when the material is dissolved in surface waters.

The size of the silica polymer can be substantial, ranging up to the colloidal state. Colloidal silica is rarely present in groundwaters. It is most commonly present in surface waters during periods of high runoff.

The polymeric form of silica does not produce color in the standard molybdate
based colorimetric test for silica. This form of silica is termed "nonreactive".
The polymeric form of silica is not thermally stable and when heated in
a boiler reverts to the basic silicate monomer, which is reactive with
molybdate.

As a result, molybdate testing of a boiler feedwater may reveal little or no silica, while boiler blowdown measurements show a level of silica that is above control limits. High boiler water silica and low feedwater values are often a first sign that colloidal silica is present in the makeup.

One method of identifying colloidal silica problems is the use of atomic emission or absorption to measure feedwater silica. This method, unlike the molybdate chemistry, measures total silica irrespective of the degree of polymerization.